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EXCHANGE 


The  Conductance   of  Aqueous  Solutions   of 

lodic  Acid  and  the  Limiting  Value 

of  the  Equivalent  Conductance 

of  the  Hydrogen  Ion 


BY 
HENRY  C  PARKER 


A  DISSERTATION 


SUBMITTED  TO   THE   FACULTY  OF  CLARK   UNIVERSITY,  WORCESTER 

MASS.,  IN  PARTIAL  FULFILMENT  OF  THE   REQUIREMENTS   FOR 

THE  DEGREE  OF  DOCTOR  OF  PHILOSOPHY,  AND  ACCEPTED 

ON  THE  RECOMMENDATION  OF  CHARLES   A.    KRAUS 


CLARK  UNIVERSITY 
1922 


EASTON,  PA. 
ESCHENBACH  PRINTING  COMPANY 


The  Conductance   of  Aqueous  Solutions   of 

lodic  Acid  and  the  Limiting  Value 

of  the  Equivalent  Conductance 

of  the  Hydrogen  Ion 


BY 
HENRY  C  PARKER 


A  DISSERTATION 


SUBMITTED  TO   THE   FACULTY  OF  CLARK   UNIVERSITY,  WORCESTER 

MASS.,  IN  PARTIAL  FULFILMENT  OF  THE   REQUIREMENTS  FOR 

THE  DEGREE  OF  DOCTOR  OF  PHILOSOPHY,  AND  ACCEPTED 

ON  THE  RECOMMENDATION  OF  CHARLES   A.    KRAUS 


CLARK  UNIVERSITY 

J922 


EASTON,  PA. 
ESCHENBACH  PRINTING  COMPANY 


v>> 


"2 

O 


THE    CONDUCTANCE    OF    AQUEOUS    SOLUTIONS    OF    IODIC 

ACID  AND  THE  LIMITING  VALUE  OF  THE  EQUIVALENT 

CONDUCTANCE  OF  THE  HYDROGEN  ION 

Introduction 

While  the  conductance  of  aqueous  solutions  of  the  salts  has  been  ex- 
tensively investigated  and  we  now  have  fairly  reliable  data  for  these  sub- 
stances to  concentrations  as  low  as  10 ~4  N  or  even  lower,  the  corresponding 
data  for  the  strong  acids  and  bases  in  aqueous  solutions  remain  very  un- 
certain. Obviously,  the  limiting  values  of  the  equivalent  conductance 
for  the  acids  and  bases  are  even  less  certain.  This  lack  of  accurate  data 
for  the  acids  and  bases  is  largely  due  to  the  fact  that  at  low  concentrations 
conductance  measurements  with  these  substances  are  attended  with 
difficulties  inherent  in  the  nature  of  the  solutions  in  question.  In  measur- 
ing the  conductance  of  a  dilute  solution  it  is  necessary,  on  the  one  hand, 
that  all  the  electrolyte  present  shall  be  in  a  known  state  and,  on  the  other, 
that  the  specific  conductance  due  to  this  electrolyte  shall  be  determinable. 
Except  in  the  case  of  salts  of  weak  acids  and  bases,  where  hydrolysis  inter- 
venes, impurities  present  in  the  water  have  little  influence  on  the  state  of  the 
salt  in  solution.  At  the  same  time,  any  impurities  present  are  not  mea- 
surably influenced  by  the  presence  of  the  salt,  and  accordingly  the  con- 
ductance due  to  these  impurities  may  be  corrected  for  by  measuring  the 
conductance  of  the  water  from  which  these  solutions  are  made  up.  With 
solutions  of  the  acids  and  bases,  this  is  not  the  case.  In  view  of  the  fact 
that  the  present  investigation  is  limited  to  an  acid,  it  is  unnecessary  to 
discuss  the  case  of  solutions  of  bases  further  than  to  state  that  the  behavior 
of  these  substances  is  similar  to  that  of  the  acids.  If  the  impurity  present 
in  the  water  from  which  an  acid  solution  is  made  up  contains  a  base,  or 
the  salt  of  a  weak  acid,  then  the  concentration  of  the  acid  will  be  influ- 
enced by  the  presence  of  this  material;  furthermore,  the  correction  to 
be  applied  for  the  conductance  due  to  the  impurity  cannot  be  determined 
by  measuring  the  conductance  of  the  original  solvent.  This  difficulty  was 
met  with  at  a  very  early  date.  It  was  found  that  the  equivalent  con- 
ductance of  an  acid,  instead  of  approaching  a  limiting  value  asymptoti- 
cally as  the  concentration  is  decreased,  passes  through  a  maximum  after 
which  it  decreases  the  more,  the  lower  the  concentration.1  It  has  been 
suggested  that  this  result  is  due  to  the  presence  of  a  base  or  a  salt  of  a 
1  Arrhenius,  Bihang  Schived.  Akad.,  8,  Nos.  13  and  14  (1884).  Kohlrausch,  Ann. 
Physik,  26,  161  (1885).  Ostwald,  /.  prakt.  Chem.,  31,  433  (1885).  Compare  also 
Kohlrausch  and  Holborn,  "Leitvermogen  der  Elektrolyte,"  Teubner,  Leipzig,  1898, 
p.  92. 


50783i 


weak  acid  in  the  solvent,  as  a  result  of  which  a  certain  amount  of  the 
hydrogen  ion  disappears  on  reaction  with  the  base  and  the  equivalent 
conductance  is  accordingly  decreased.2  This  decrease  will,  of  course, 
be  relatively  the  greater,  the  lower  the  concentration  of  the  acid.  Thus 
far,  however,  no  measurements  have  been  carried  out  in  which  this  source 
of  error  has  been  eliminated. 

Indeed,  the  error  in  the  case  of  the  acids  may  be  traced  to  two  sources: 
first,  impurities  present  in  the  water  used  in  making  up  the  solution;  and 
second,  impurities  dissolved  from  the  containing  vessels.  The  first  of 
these  sources  of  error  has  been  recognized  as  mentioned  above,  but  the 
second  source  of  error  appears  not  to  have  been  considered  heretofore. 
It  is  well  known,  however,  that  even  the  best  of  glass  is  appreciably  soluble 
in  acids  and  it  is  accordingly  to  be  expected  that  experiments  carried 
out  in  glass  cells  will  be  measurably  affected  at  low  concentrations  owing 
to  the  solubility  of  the  glass.  In  order  to  overcome  these  difficulties, 
therefore,  it  is  necessary  to  carry  out  measurements  with  water  which 
has  been  freed  from  all  impurities  and  in  cells  which  are  resistant  to  the 
action  of  acids.  In  the  present  investigation,  pure  water  has  been  pre- 
pared by  a  method  previously  developed  in  this  Laboratory3  and  measure- 
ments have  been  carried  out  in  a  quartz  cell. 

In  selecting  an  acid  for  the  purpose  of  this  investigation,  it  appeared 
desirable,  if  possible,  that  the  acid  should  be  relatively  strong  in  order  to 
reduce  the  errors  due  to  extrapolation  and  that  it  should  be  readily  manip- 
ulated for  the  purpose  of  making  up  the  solutions.  For  this  reason  iodic 
acid  was  chosen,  since  it  is  readily  prepared  in  a  pure  state  as  a  solid, 
in  which  condition  it  may  be  weighed  and  transferred  to  the  cell,  while 
at  the  same  time  it  is  but  little  weaker  than  hydrochloric  acid. 

Preparation  of  Materials 

Iodic  Acid. — The  iodic  acid  employed  in  this  investigation  was  prepared  from  the 
best  quality  of  iodic  acid  obtainable  on  the  market.  This  acid  was  purified  by  repeated 
recrystallization.  Considerable  difficulty  was  at  first  experienced  in  the  process  of 
recrystallization  owing  to  the  fa^ct  that,  on  evaporating  an  aqueous  solution,  the  acid 
forms  a  supersaturated  solution  which  is  brought  to  crystallization  with  difficulty,  even 
on  inoculation.  It  was  found,  however,  that  when  a  small  amount  of  nitric  acid  is 
added  to  the  water  solution  and  the  water  is  extracted  slowly  in  a  desiccator,  beautiful 
crystals  of  the  acid  are  formed.  A  beaker  containing  sodium  hydroxide  was  placed  in  a 
desiccator  with  sulfuric  acid,  together  with  the  solution  of  iodic  acid.  The  purpose  of  the 
sodium  hydroxide  was  to  prevent  the  concentration  of  the  nitric  acid  from  reaching  too 
high  a  value.  After  precipitation,  the  crystals  of  iodic  acid  were  washed  with  very  dilute 
nitric  acid  and  then  dried  over  a  solution  of  potassium  hydroxide.  This  was  for  the 
purpose  of  removing  any  possible  remaining  traces  of  nitric  acid.  The  last  traces  of 
moisture  were  removed  over  sulfuric  acid.  The  final  process  of  desiccation  was  carried 

1  Whetham  and  Paine,  Proc.  Roy.  Soc.,  81A,  58  (1908).  Paine  and  Evans,  Proc. 
Cambridge  Phil.  Soc.,  18,  1  (1914). 

*  Dexter,  Dissertation,  Clark  University,  1917.     J.  Am.  Chem.  Soc.,  44,  2468  (1922). 


5 

out  shortly  before  the  iodic  acid  was  to  be  used,  since  otherwise,  as  was  found,  a  certain 
amount  of  the  anhydride,  I^Os,  is  liable  to  be  formed.  The  process  of  recrystallization 
above  described  was  carried  out  4  times. 

The  purity  of  the  acid  was  tested  by  various  methods.  First,  the  acidity  was  de- 
termined by  comparison  with  a  solution  of  hydrochloric  acid  which  had  been  standard- 
ized against  silver  by  weighing  as  silver  chloride.  The  average  of  three  determinations 
gave  for  the  normality  of  the  acid  the  value  0. 10273  =*=  3  N.  The  sodiumhy  droxide 
solution  used  compared  with  this  standard  acid,  in  two  titrations,  gave  the  factor 
0.099193  =*=  11,  weight  burets  being  employed.  On  titrating  the  iodic  acid  against  this 
solution  of  sodium  hydroxide,  the  results  given  in  Table  I  were  obtained. 

TABLE  I 

DETERMINATION  OF  ACIDITY  OP  IODIC  Aero 

HIO3  taken  NaOH  HIO3  found  Error 

G.  G.  G.  % 

0.74285  42.681  0.74307  +0.03 

0.51119  39.318  0.51162  +0.08 

0.85278  48.9185  0.85367  +0.016 

The  acidity  was  likewise  determined  by  a  new  method  which  consists  in  precipitating 
the  iodate  with  silver  as  silver  iodate  and  weighing.  Since  silver  iodate  is  somewhat 
soluble,  the  filtrate  is  saved  and  analyzed,  iodine  being  liberated  with  potassium  iodide 
and  sulfuric  acid  and  titrated  against  standard  thiosulfate  solutions.  The  results  of 
this  method  are  given  in  Table  II. 

TABLE  II 

DETERMINATION  OF  IODIC  ACID  AS  SILVER  IODATE 

HIO3  taken            AgIO3     HIO3  in  filtrate  Total  HIO3  found  Error 

G.                       G.                     G.  G.  % 

0.21224             0.33721            0.0027  0.21247  +0.04 

0.15159             0.24019            0.00226  0.15168  +0.07 

0.16769             0.26617            0.00192  0.16750  -0.11 

The  nitrate  content  was  shown  to  be  negligible  by  a  colorimetric  test,  using  phenol- 
sulfonic  acid,  the  iodate  being  precipitated  with  silver  sulfate.  This  test  showed  roughly 
6  parts  of  nitrogen  as  nitrate  per  million  of  iodate.  An  attempt  was  also  made  to  deter- 
mine the  iodic  acid  by  dehydrating  to  anhydride  and  weighing  the  difference,  as  well  as  by 
liberating  the  iodine  by  means  of  potassium  iodide  and  estimating  by  means  of  standard 
thiosulfate  solution.  Neither  of  these  latter  tests  proved  entirely  satisfactory,  however, 
the  former  on  account  of  the  difficulty  of  perfect  dehydration,  and  the  latter  owing  to 
the  large  amount  of  iodine  liberated.  The  former  tests,  however,  in  ail  cases  fell  within 
the  limit  of  experimental  error  and  were  thus  satisfactory  from  that  point  of  view. 

The  best  proof  of  the  purity  of  the  iodic  acid,  however,  consists  in  the  agreement  of 
the  values  obtained  with  samples  of  iodic  acid  from  the  different  crystallizations.  The 
precision  of  the  conductance  measurements  is  far  greater  than  that  of  any  analytical 
method  which  is  here  applicable.  The  four  runs  in  dilute  solutions  given  below  repre- 
sent results  with  three  different  crystallizations  of  the  acid ;  and,  as  may  be  seen,  these 
agree  very  closely. 

Water. — The  water  employed  in  this  investigation  was  purified  by  the  method  de- 
veloped in  this  Laboratory.  Ordinary  water  was  distilled  from  an  alkaline 
permanganate  solution  under  the  usual  conditions,  the  first  fraction  of  the  distillation 
being  discarded.  The  product  obtained  was  then  again  distilled  from  a  dil.  alkaline 
permanganate  solution  in  a  special  still  of  the  type  mentioned  above  in  which  carbon 


dioxide  was  removed  by  fractional  condensation.  In  order  to  avoid  contamination  due 
to  the  surrounding  atmosphere,  distillation  was  carried  out  under  a  slight  pressure  of  air 
which  had  previously  been  purified,  as  will  be  described  below.  The  water  employed  in 
these  measurements  had  a  specific  conductance  varying  between 
0.09  X  10  ~9  and  0.5  X  10  ~6.  As  will  be  seen  below,  with  water 
having  the  higher  value  of  the  specific  conductance,  the  influence 
of  the  impurities  could  be  detected  in  the  conductance  values. 
Purification  of  Air. — Pure  air  was  necessary  for  the  purpose 
of  stirring  the  solution  in  the  cell  as  well  as  for  the  preparation 
of  the  water  as  stated  above.  The  air  used  for  this  purpose  was 
purified  by  means  of  a  special  set  of  continuously  acting  towers, 
a  sketch  of  which  is  shown  in  Fig.  1.  The  air  enters  through 
the  tube  A,  carrying  before  it  a  column  of  the  purifying  solu- 
tion which  is  raised  to  the  top  of  the  tower  and  there  projected 
over  the  beads  with  which  the  tower  is  filled.  The  air  then 
passes  down  among  the  beads  over  which  the  solution  is  con- 
tinuously flowing.  The  chamber  containing  the  beads  has  a 
length  of  approximately  50  cm.  and  a  diameter  of  3  cm.  so  that 
the  rate  of  flow  of  the  air  through  the  purifying  apparatus  is 
relatively  low.  This  method  of  purification  was  found  very 
successful.  In  order  to  remove  both  ammonia  and  carbon 
dioxide,  3  towers  were  employed,  the  first  of  which  was  filled 
with  a  solution  of  sulfuric  acid  and  the  remaining  2  with  solu- 
tions of  sodium  hydroxide.  The  necessary  pressure  for  the  air 
was  obtained  by  means  of  an  aspirator. 


Fig.  1 .  — Construction 
of  air  purifying 
towers. 


Measuring  Apparatus 

Bridge. — A  drum-wound,  slide-wire  bridge  with  extension  coils  was  used  in  measur- 
ing the  resistance  of  the  solutions.  The  bridge  wire  was  calibrated  by  the  Kelvin  method, 
and  by  means  of  the  corrections  obtained  it  was  possible  to  check  resistance  readings  to 
better  than  0.01%.  Air  condensers  were  employed  to  balance  out  the  effects  of  induc- 
tance and  capacity  in  the  circuit.  The  entire  apparatus  was  carefully  shielded,  all  con- 
nections being  lead  covered  and  grounded  and  all  measuring  instruments  being  kept  with- 
in a  lead-lined  case. 

Source  of  E.M.F. — A  Vreeland  oscillator  was  employed  as  a  source  of  alternating 
current.     The  oscillator  was  fitted  out  with  a  variable  air  transformer  by  means  of  which 
it  was  possible  to  regulate  the  voltage,  while  the  frequency  of  the  oscillator  was  adjusted 
by  means  of  suitable  capacities  introduced  in 
the  oscillator  circuit.     In  general,  4  combina- 
tions of  capacities  were  employed  and  the  fre- 
quencies corresponding  to  these  were  deter- 
mined by  means  of  an  oscillograph.     Prints  of 
the  records  thus  obtained  are  shown  in  Fig.  2. 
The  horizontal  dashes  appearing  in  the  figure 
are  due  to  the  timing  device.     It  is  seen  that 
the  oscillator  gave  a  perfect  sine  wave.     Since 


Fig.  2. — Oscillogram  of  Vreeland 
oscillator. 


the  oscillograph  had  a  frequency  considerably  above  10,000,  any  overtones  which  might 
have  been  present  would  unquestionably  have  appeared  on  the  plate.4     The  four  fre- 

4  This  result  would  appear  to  contradict  the  statement  of  Eyster   [/.  Am.  Inst. 
Elec.  Eng.,  39,  889  (1920)]  to  the  effect  that  the  Vreeland  oscillator  does  not  give  a  sine 


quencies  thus  determined  were  found  to  be  1747,  1217,  986.5  and  489.5  per  second.  It 
was  found  that  in  most  cases  a  frequency  of  1217  gave  the  most  satisfactory  result,  both 
as  regards  the  accuracy  of  the  readings  and  the  distinctness  of  the  setting.  This  fre- 
quency was,  therefore,  used  practically  throughout  this  work. 

Resistances. — The  resistance  boxes  consisted  of  Curtis-wound  coils,  ranging  from 
10  to  10,000  ohms  capacity.  These  were  calibrated  against  resistance  standards  which 
had  recently  been  calibrated  at  the  Bureau  of  Standards.  In  the  older  measurements, 
which  were  carried  out  with  solutions  contained  in  glass  cells,  the  resistances  had  not 
been  calibrated,  for  which  reason  the  measurements  are  not  as  accurate  as  those  later 
obtained.  These  earlier  results,  however,  have  not  been  used  for  the  purpose  of  deter- 
mining the  final  conductance  values.  In  the  later  measurements  with  quartz  and  Pyrex 
glass  cells,  which  were  begun  in  about  September,  1919,  and  completed  in  June,  1920, 
calibration  corrections  were  made.  Two  series  of  calibrations  of  the  resistances  were 
carried  out,  one  in  September,  1919,  and  another  in  May,  1921.  All  the  coils  below  10,000 
ohms  were  found  to  be  constant  over  this  interval  to  better  than  0.01%,  but  the  10,000 
ohm  coils  showed  a  variation  of  0.88%.  Accordingly,  in  correcting  the  measurements, 
it  was  assumed  that  the  resistance  varied  as  a  linear  function  of  the  time,  and  a  correction 
applied  which  amounted  to  0.044%  per  month.  Four  standard  resistances  of  10,  100, 
1,000  and  10,000  ohms  were  employed.  These  were  found  to  be  mutually  consistent. 
All  calibrations  were  carried  out  with  a  direct  current  and  galvanometer.  With  all 
corrections  applied,  it  was  found  possible  to  obtain  check  measurements  within  0.005%. 

The  Measuring  Cells. — In  all,  5  cells  were  used  in  measuring  the  conductance  of 
the  acid,  in  addition  to  an  auxiliary  cell  of  the  pipet  form  which  was  employed  for  the 


Fig.  3. '— Construction  of  various  types  of  cells  employed. 

purpose  of  calibrating  the  measuring  cells.  Two  of  the  measuring  cells  were  constructed 
of  soda-lime  glass,  and  two  of  quartz,  and  a  fifth  cell  was  constructed  of  Pyrex  glass. 
The  cells  are  shown  in  outline  hi  Fig.  3.  In  this  figure,  V  A  and  V  B  represent  front  and 
side  elevations,  respectively,  of  the  glass  cell  which  was  employed  in  measuring  the  con- 
ductance of  the  dilute  solutions.  The  electrodes  of  this  cell  had  an  area  of  approximately 
1.5  sq.  cm.  and  were  placed  4  cm.  apart.  The  body  of  the  cell  was  approximately  60 


cm.  long  and  5  cm.  in  diameter  and  had  a  volume  of  approximately  1  liter.  The  glass 
cell  used  for  measurements  at  higher  concentrations  had  electrodes  of  the  same  cross 
section  as  Cell  V  which  were  placed  at  a  distance  of  17  cm.  from  each  other.  This  cell 
is  shown  as  VI  A  and  VI  B  in  the  figure.  The  height  of  this  cell  was  approximately  30  cm. 
and  its  volume  approximately  200  cc.  The  electrodes  in  all  cells  were  platinized  accord- 
ing to  the  method  of  Kohlrausch,  and  were  then  heated  in  a  blowpipe  flame  until  gray. 

In  the  same  figure,  II  represents  the  large  quartz  cell  employed  in  measurements 
with  the  more  dilute  solutions.  This  cell  consisted  of  a  Vitreosil  flask  having  a  volume  a 
little  over  3  liters.  The  electrodes  had  an  area  of  1.5  sq.  cm.  and  were  placed  at  a  dis- 
tance of  2.5  cm.  from  each  other.  The  electrodes  were  provided  with  a  brace  under- 
neath in  order  to  guard  against  displacement.  The  electrode  stems  were  constructed 
of  glass,  but  the  tube  F,  through  which  air  was  blown  for  the  purpose  of  stirring  the  solu- 
tion, was  constructed  of  quartz.  A  rubber  stopper  D  served  to  hold  the  electrodes  in 
position  in  the  cell.  The  cell. was  calibrated  in  order  to  determine  the  effect  of  the  posi- 
tion of  the  electrodes,  especially  with  reference  to  the  level  of  the  liquid.  A  curve  of 
resistance  readings  for  different  heights  of  the  liquid  was  made  and  the  height  corre- 
sponding to  a  minimum  distance  was  chosen  for  the  position  of  the  electrodes.  In  this 
position  a  difference  of  1.25  cm.  in  the  level  of  the  liquid  was  required  to  make  a  difference 
of  0.01%  in  the  resistance.  The  quartz  cell  used  for  carrying  out  the  measurements  at 
higher  concentrations  is  shown  as  IV  A  and  IV  B .  Circular  electrodes,  having  an  area  of 
approximately  1.5  sq.  cm.,  were  placed  in  a  horizontal  quartz  tube  of  30  cm.  length  and 
of  2  cm.  diameter.  The  leads  were  introduced  after  the  manner  described  by  Kraus.6 
The  form  of  the  cell  as  here  shown  was  so  designed  that  its  contents  could  readily  be 
mixed  by  shaking  without  removing  it  from  the  thermostat.  The  capacity  of  this  cell 
was  approximately  1.5  liters.  The  construction  of  the  Pyrex  glass  cell  was  similar  to  that 
of  Cell  IV. 

The  Thermostat. — The  thermostat  employed  in  these  measurements  was  kept  at 
a  constant  temperature  within  0.002°  by  means  of  a  steel-contained  mercury  thermo- 
regulator.  In  carrying  out  the  measurements  with  the  glass  cells,  the  thermostat  was 
filled  with  water,  but  in  all  the  later  measurements  the  thermostat  was  filled  with  kero- 
sene. At  very  low  concentrations,  where  the  resistance  of  the  cell  is  large,  capacity 
effects  apparently  are  introduced  when  a  water-filled  thermostat  is  employed,  and  it  was 
for  this  reason  that  kerosene  was  later  employed  in  place  of  water. 

Balance. — The  balance  used  for  weighing  out  the  acid  was  a  standard  analytical 
balance  especially  adjusted  for  a  sensitivity  of  0.02  mg.  It  was  found  possible  to  obtain 
check  weighings  to  0.01  mg.  The  weights  used  were  calibrated  against  a  set  of  assay 
weights  which  had  been  standardized  by  the  Bureau  of  Standards. 

Preliminary  Measurements 

Cell  Constant. — The  values  of  the  cell  constant  throughout  this 
investigation  are  based  upon  the  measurements  of  Kohlrausch  and  Maltby6 
with  solutions  of  potassium  chloride  at  18°.  The  constant  of  the  auxiliary 
cell,  from  which  the  constants  of  the  other  cells  were  determined  by  inter- 
comparison,  was  determined  at  18°,  according  to  the  method  described 
in  the  preceding  paper  by  the  present  authors. 

The  cell  constants  as  determined  with  a  series  of  independent  solutions 
are  given  in  Tables  III  and  IV.  In  Table  III  are  given  values  of  the  con- 

*  Kraus,  U.  S.  pat.,  No.  1,093,997,  1914. 

8  Kohlrausch  and  Maltby,  Wiss.  Abh.  Phys.-Tech.  ReichsansL,  3,  157  (1900). 


9 

stant  for  the  standard  cell  I  as  determined  at  different  times  during  the 
course  of  the  investigation. 

TABUS  III 

CONSTANT  OF  STANDARD  CELL  I 

Before  Run  I  After  Run  V  After  Run  VII 

3.67125  3.67039  3.67041 

3.67113  3.67098  3.67037 

3.67108 
3.67123 
Av.  3.67117  3.67068  3.67039 

It  will  be  seen  that  there  is  a  slow  drift  of  these  values  with  the  time.  A 
sufficient  check,  however,  was  kept  upon  this  variation  so  that  the  cell 
constant  as  determined  for  the  measuring  cells  may  be  relied  upon.  In 
Table  IV  are  given  the  constants  of  the  large  quartz  cell  IT.  the  Pyrex 
glass  cell  III,  the  small  quartz  cell  IV,  the  large  lime  glass  cell  V,  and  the 
small  lime  glass  cell  VI. 

TABLE  IV 
CONSTANTS  OF  DIFFERENT  CELLS  USED  IN  MEASURING  THE  CONDUCTANCE  OP  TODIC 

Acm 
Cell  II  Cell  III 


Before 
Run  I 

Before 
Run  II 

Before 
Runs  III,  IV,  V 

After 
RunV 

0.28616 

0.285436 

0.284748 

0.284697 

8.26770 

0.28644 

0.285379 

0.284711 

0.284796 

8.26665 

0.285391 

0.284741 

Av.   0.28630  0.285402  0.284733  0.284747  8.26717 

Cell  IV  Cell  V  Cell  VI 


Run  VI 

Run  VII 

6.98116 

6.99346 

0.715824 

6.35052 

6.98186 

6.99317 

0.715403 

6,35013 

6.99278 

Av.   6.98151  6.99313  0.715613  6.35033 

The  constants  of  Cells  III,  IV  and  VI  were  determined  at  25°,  using 
the  specific  conductance  given  in  the  preceding  article.  The  values  of 
the  constant  of  Cell  II,  which  was  used  for  the  most  accurate  determina- 
tions in  dilute  solutions,  are  given  as  determined  at  different  times.  These 
determinations  were  made  each  time  with  a  series  of  independent  solutions 
made  up  in  the  cell  and  then  inter  compared  with  the  standard  Cell  I. 
The  resistance  of  the  solution  in  Cell  II  was  usually  in  the  neighborhood 
of  400  ohms,  although  at  one  time  the  resistance  of  the  solution  was  varied 
between  250  and  1300  ohms  in  order  to  determine,  if  possible,  whether 
the  cell  constant  varies  as  a  function  of  the  concentration.  The  measure- 
ments in  the  more  dilute  solutions,  however,  did  not  check  well  because 


10 

of  various  sources  of  error.7    The  cell  constants  finally  employed  in  this 
investigation  are  the  mean  values  given  at  the  bottom  of  Tables  III  and 

IV. 

The  actual  measurements  on  iodic  acid,  and  most  of  the  measurements 
on  the  cell  constants,  were  carried  out  at  a  temperature  of  24.958°  instead 
of  25°.  All  data  have  been  converted  to  the  corrected  temperature  by 
factors  which  were  determined  by  measurements  made  at  the  two  tem- 
peratures with  a  series  of  solutions.  Thsee  factors  are  as  follows:  to 
convert  the  resistance  of  the  potassium  chloride  solution  from  24.958°  to 
25°,  multiply  by  0.999202;  to  convert  the  resistance  of  iodic  acid  solu- 
tions from  24.958°  to  25°,  multiply  by  0.999434.  In  determining  these 
factors,  the  solutions  were  carried  back  and  forth  between  the  tempera- 
tures in  question  several  times.  The  readings  agreed  to  better  than  0.01  %. 
The  values  given  are  the  averages  of  all  the  readings  taken. 

Temperature. — The  temperatures  of  18°  and  25°  were  established 
by  means  of  2  Beckmann  thermometers  which  had  been  calibrated  against 
a  platinum  resistance  thermometer.8  A  series -of  comparisons  were  made 
over  a  range  of  temperatures  in  the  neighborhood  of  the  temperatures  in 
question  and  the  results  were  plotted  in  order  to  determine  the  readings 
at  18°  and  25°.  The  temperature  interval  from  18°  to  25°,  which  had 
been  used  previous  to  this  calibration,  was  found  to  be  correct  to  0.01°. 
The  final  calibration  was  found  to  agree  exactly  with  the  reading  of  a 
mercury  thermometer  which  had  been  calibrated  by  the  Reichsanstalt 
in  1910. 

Density  Measurements. — In  order  to  reduce  the  concentrations  to 
a  volume  normal  basis,  it  was  necessary  to  determine  the  density  of  iodic 
acid  solutions.  The  results  of  these  measurements  are  given  in  Table  V 
and  are  shown  graphically  in  Fig.  4.  The  results  of  Heydweiller  and 
Groschuff,  made  at  18°  and  0°,  respectively,  and  reduced  to  25°,  are  also 
shown  on  the  figure.  In  correcting  these  results  to  25°,  it  was  assumed 
that  the  density  change  of  the  solution  between  the  temperatures  in  ques- 
tion is  the  same  as  that  of  pure  water  between  the  same  temperatures. 
As  may  be  seen  from  the  figure,  the  three  series  of  measurements  check 
closely  up  to  a  concentration  of  about  0.5  N.  It  is  probable  that  above 

7  A  further  investigation,  in  order  to  determine  whether  or  not  the  cell  constant 
varies  as  a  function  of  the  concentration,  has  since  been  carried   out  by   the  author 
using  a  very  accurate  method  of  intercomparison.     This  investigation  showed  that  the 
constant  of  this  cell  and,  in  fact,  the  constant  of  every  cell  investigated,   varies  appre- 
ciably with  the  concentration.     The  results  will  appear  shortly  in  a  publication  from  this 
Laboratory. 

8  The  thermometer  in  question  is  one  which  was  in  use  in  the  Research  Laboratory 
of  Physical  Chemistry  of  the  Massachusetts  Institute  of  Technology  and  which  had 
been  carefully  calibrated  by  Dr.  James  A.  Beattie.      The  author  wishes  to  express  his  in- 
debtedness to    Dr.  Beattie  for  his  assistance  in  connection  with  the  calibration  and 
also  to  Dr.  Frederick  G.  Keyes,  Director  of  the  Research  Laboratory. 


11 


this  concentration  the  assumption  made  for  the  reduction  of  the  results 
of  Heydweiller  and  Groschuff  to  25°  does  not  hold.     For  the  measure- 


;x 


1.16 

y 

.12 

/ 

x 

^x^ 

a 

x 

j] 

X 

,O4 

A 

DENSITY  AT     i5'             =   O 

<S 

HIYDWEIUER                     =   C 

j? 

GROSC.HIFF                        =   -t 

1.00 

OQft 

^          \ 

0 


0.2 


0.8 


1.0 


0.4  0.6 

Concentration. 
Fig.  4. — Density  of  iodic  acid  solutions  at  25°. 

ments  at  0.08373  N  and  0.68879  N  a  pycnometer  was  used,  while  at  the 
other  concentrations  given  in  the  table  the  results  were  obtained  by 
means  of  a  Mohr's  balance.  The  final  values  employed  were  read  from 
the  smooth  curve  shown  in  the  figure. 


Cone. 
0.08373 
0 . 144854 
0.318609 
0.40788 


TABLE  V 
DENSITY  OF  IODIC  ACID  SOLUTIONS  AT  25° 

df  Cone. 

1.00902  0.66427    - 

1.01878  0.68879 

1.04379  0.91111 
1.05675 


j25 

d4 

1.09467 
1.09768 
1.12957 


Resistance  of  Leads. — The  resistance  of  the  leads  to  the  cells  was  in 
each  case  determined  by  filling  the  cell  with  mercury  and  measuring  the 
resistance  against  a  series  of  low  resistances,  using  direct  current  and  a 
galvanometer.  The  resistance  of  the  wires  leading  from  the  bridge  to  the 
resistance  box  had  to  be  taken  into  consideration.  In  view  of  the  fact 
that  readings  were  carried  out  with  the  bridge  setting  near  the  middle 
of  the  bridge,  sufficient  precision  was  obtained  by  subtracting  from  the 
total  calculated  resistance  the  difference  between  the  resistance  of  the 
leads  in  the  cell  and  the  remaining  leads  in  the  bridge  set-up.  The  values 
of  these  differences  as  determined  were  as  follows:  Cell  I,  0.047  ohms; 
Cell  II,  0.032  ohms;  Cell  III,  0.020  ohms;  Cell  IV,  0.044  ohms;  Cell  V, 
0.083  ohms;  Cell  VI,  0.092  ohms. 


12 

Variation  of  Cell  Resistance  with  Frequency  and  Potential.— Taylor 
and  Acree9,  in  their  investigation,  found  a  variation  of  the  cell  resistance 
with  the  frequency  of  the  current  and  also  with  the  value  of  the  impressed 
voltage.  These  effects  were  also  observed  in  the  present  investigation. 
It  was  always  found  possible  to  eliminate  the  change  of  resistance  with 
the  frequency  by  sufficient  platinization  of  the  electrodes  so  that  the  dif- 
ference in  resistance  between  500  and  1700  cycles  was  as  low  as  0.01%. 
In  measurements  with  the  quartz  cell,  no  readings  were  made  with  re- 
sistances below  300  ohms,  since  the  above-mentioned  effect  was  found  to 
increase  greatly  at  low  resistances.  The  change  of  resistance  with  the 
voltage  was  more  difficult  to  eliminate.  It  was  suggested  by  Taylor  and 
Acree  that  this  change  is  due  to  contamination  of  the  solution  in  the  cell, 
but  it  is  difficult  to  believe  that  this  is  the  sole  reason  for  this  behavior, 
since  the  effect  was  observed  in  the  quartz  cell  after  one  or  two  buckets — 
scrupulously  clean — had  been  dropped  into  conductivity  water  of  a  spe- 
cific conductance  of  only  0.09  X  10~6.  A  voltage  change  from  0.2  to  7 
volts  usually  produced  a  change  in  the  apparent  resistance  of  about  0.02%. 
In  the  case  of  the  iodic  acid  solutions,  it  was  noticed  that  the  higher  voltage 
gave  the  lower  resistance,  while  in  that  of  potassium  chloride  solutions  the 
effect  was  in  the  opposite  direction  and  of  about  the  same  order  of  magni- 
tude. This  effect  was  also  observed  when  film  resistances  were  employed 
in  place  of  the  Curtis  coils,  thus  showing  that  the  phenomenon  is  localized 
in  the  cell  itself.  It  was  not  found  possible  to  locate  the  cause  of  this 
effect  in  the  limited  time  available  and,  since  it  was  necessary  to  use  a 
higher  voltage  in  order  to  obtain  reliable  bridge  readings  at  the  higher 
resistances,  it  was  decided  to  employ  the  higher  voltage  throughout  the 
measurements.  With  low  resistances,  the  higher  voltage  caused  some 
difficulty  as  a  result  of  thermal  effects  but  this  was  eliminated  by  rapid 
readings. 

Experimental  Procedure  and  Results 

Manipulation. — The  method  of  carrying  out  a  series  of  conductance 
measurements  was  as  follows.  The  iodic  acid  was  weighed  in  small  quartz 
buckets  which  were  hung  on  a  platinum  wire  in  a  desiccator.  These 
buckets  were  all  weighed  and  checked  shortly  before  making  a  run.  In 
the  case  of  the  runs  carried  out  with  the  quartz  cells,  the  buckets  were 
provided  with  a  small  bulb  at  the  top  which  served  to  float  them,  thus 
keeping  them  out  of  the  way  of  the  electrodes  and  also  providing  a  quick 
and  convenient  method  of  obtaining  concentration  equilibrium  in  the 
solution.  At  the  bottom,  the  buckets  were  provided  with  a  small  opening 
through  which  the  solution  flowed  out  as  rapidly  as  formed.  With  the 
use  of  these  buckets,  it  was  usually  possible  to  obtain  concentration  equi- 
librium at  the  end  of  a  period  of  10  minutes.  The  buckets  were  cleaned 
9  Taylor  and  Acree,  /.  Am.  Chem.  Soc.,  38,  2415  (1916). 


13 

by  treating  with  cleaning  mixture,  boiling  water,  and  finally  by  steaming 
and  drying  over  an  electric  hot  plate.  The  cells  were  cleaned  inside  and 
out  with  hot  cleaning  mixture,  after  which  hot  water  was  allowed  to  run 
through  them  for  an  hour.  They  were  then  steamed,  with  frequent  rinsing, 
and  finally  they  were  attached  to  the  still.  In  the  case  of  the  quartz  cells, 
a  slight  pressure  of  purified  air  was  applied  in  filling  in  order  to  prevent 

TABLE  VI 
CONDUCTANCE  OF  IODIC  ACID  IN  QUARTZ 


Run  1,  Cell  II, 

K  =  0.28630, 

M  =  0.5  X  10-«, 

W  =  3127.70 

Cone.  X  103 

A 

Cone.  X  10» 

A 

0.0762746 

386.308 

0.247254 

387.218 

0  .  148601 

386.875 

0.426920 

387.061 

Run  2,  Cell  II, 

K  =  0.285402, 

M  =  0.115  X  10~6, 

W  =  3187.98 

0.0537214 

388.673 

0.410722 

387.496 

0.131132 

388.799 

0.681964 

386.035 

0.232533 

388.295 

1  .  10967 

384.076 

Run  3,  Cell  III, 

K  =  0.284733, 

//  =  0.0976  X  10~6, 

W  =  3169.38 

0.0664405 

389.046 

0.439514 

387.207 

0.142177 

388.535 

0.703408 

385.963 

0.251694 

388.192 

1.10330 

384.217 

Run  4,  Cell  II, 

K  =  0.284733, 

M  =  0.095  X  10-«, 

W  =  3206.57 

0.0685165 

389.018 

0.648686 

386.117 

0.144572 

388.940 

0.986010 

384.543 

0.233308 

388.311 

1  .49429 

382.360 

0.399158 

387.462 

2.11108 

380.059 

Run  5,  Cell  II, 

K  =  0.284733, 

ju  =  0.0989  X  10-«, 

W  =  3226.16 

0.0962906 

389.022 

0.710642 

385.764 

0  .  168284 

388.616 

1.02335 

384.310 

0.286320 

388.036 

1.52809 

382.259 

0.465172 

386.999 

2.10323 

380.092 

Run  6,  Cell  IV, 

K  =  6.98151, 

M  =  0.46  X  10~6, 

W  =  993.332 

0.623489 

384.974 

9.03267 

362.151 

1.73067 

380.918 

16.1179 

349.421 

3.07736 

376.582 

30.5179 

330.179 

5.38453 

370.333 

Run  7,  Cell  IV, 

K  =  6.99313, 

M  =  0.38  X  10~6, 

W  =  951.129 

0.00143615 

382  .  195 

0.0183879 

345.876 

0.00304044 

376.774 

0.0349757 

325.852 

0.00547988 

370.130 

0.0673634 

298.053 

0.00969682 

360.794 

0.193258 

244.359 

contamination  from  the  external  atmosphere.  In  collecting  the  water 
in  the  cell,  steam  was  first  allowed  to  blow  through  the  cells  for  a  time, 
during  which,  also,  they  were  thoroughly  rinsed  with  water  very 
nearly  at  its  boiling  point.  Finally,  the  condensed  water  was  cooled  by 
means  of  a  suitable  cooler  and  the  cell  was  repeatedly  rinsed  out  by  filling 
and  siphoning  off  the  water.  Throughout  these  operations,  the  resistance 


14 

of  the  water  in  the  cell  was  measured  and  when  water  of  a  desired  quality 
was  obtained  the  cell  was  filled. 

After  filling,  the  cell  was  removed  from  the  still,  weighed,  and  placed 
in  a  thermostat  where  it  was  left  for  3  hours  to  come  to  temperature, 
this  being  the  longest  time  necessary  in  order  to  insure  equilibrium,  ac- 
cording to  a  series  of  measurements.  When  temperature  equilibrium 
was  established,  the  first  quartz  bucket  was  introduced  and  resistance 
measurements  were  made  every  15  minutes.  In  the  case  of  the  first 

TABUS  VII 

CONDUCTANCE  OP  IODIC  ACID  IN  LIME  GLASS  CELLS 
Run  8,  Cell  V,  K  =  0.715613,          n  =  0.8  X  10~6,       W  =  917.867 

Cone.  X  10»  A  Cone.  X  10"  A 

0.17619  376.313  1.4305  '      381.416 

0.36624  380.412  2.8999  376.974 

0.54857  382.745  4.3311  373.094 

0.70948  382.775  5.7854  369.541 

1.0938  382.316 

Run  9,  Cell  V,  K  =  0.715613,  n  =  0.24  X  lO"8,       W  =  892.875 

0.096707  373.45  1.3786  381.634 

0.241261  381.879  2.0890  379.363 

0.37173  383.475  2.7931  377.244 

0.58002  383.684  4.2595  373.107 

0.83381  383.362  8.5392  363.322 

1.1043  382.413 

Run  10,  Cell  VI,  K  =  6.35033,  /z  =  1 .05  X  10 ~8,  W  =  103.504 

0.0041409  373.509  0.066075  298.911 

0.019662  344.082  0.087229  285.669 

0.025734  335.869  0.12811  266.024 

0.036047  324.093  0.25030  229.872 

0.046199  314.358  0.49329  191.609 

bucket,  with  Cell  II,  it  was  found  necessary  to  assume  that  the  minimum 
resistance  observed  was  the  correct  one,  since  there  was  a  gradual  change 
with  the  time  extending  over  a  period  of  several  hours,  due  to  the  solution 
of  glass  from  the  leads  carrying  the  electrodes.  When  the  resistance 
began  thus  to  increase  the  second  bucket  was  introduced,  and  from  then 
on  resistance  measurements  were  continued  at  each  point  until  no  drift 
was  apparent  in  the  readings.  The  effect  in  question  was  negligible  in  all 
cases  in  which  the  contents  of  the  bucket  dissolved  promptly. 

Conductance  Data. — The  results  obtained  in  this  investigation  are  given 
in  Tables  VI,  VII  and  VIII.  At  the  head  of  each  sub-table  is  given  a 
number  indicating  the  cell  used,  the  constant  of  the  cell  K,  the  specific 
conductance  of  the  water  employed,  //,  and  the  total  weight,  W,  of  the 
water  employed  in  making  up  the  solution. 

All  weights  given  are  reduced  to  a  vacuum.  The  value  of  the  molecular 
weight  of  iodic  acid  was  assumed  to  be  175.928. 


15 


TABLE  VIII 

CONDUCTANCE  OP  IODIC  Aero  IN  PYREX  GLASS  CELL 
Run  11,  Cell  III,  K  =  8.26717,  M  =.0.53  X  10-«,     W  =  367.578 

Cone.  X  10»  A 

1-  63835  381.261 

3.62966  374.952 

The  results  are  shown  graphically  in  Figs.  5  and  6.  In  Fig.  5,  values 
of  A  are  plotted  against  values  of  the  logarithm  of  the  concentration, 
while  in  Fig.  6  values  of  I/A  are  plotted  against  values  of  the  specific 
conductance. 


390 


365 


4.2 


4.4 


4.6 


3.4 


3.6 


3.8 


4.8  3.0  3.2 

Log  concentration. 

Fig.   5. — Influence   of   impurities    on  the  conductance  of  dilute  solutions    of 

iodic  acid. 

From  an  examination  of  Fig.  6,  it  will  be  seen  that  in  Run  1  the  curve 
exhibits  a  minimum  point  at  a  concentration  of  approximately  1.3  X  10~4 
Nt  after  which  it  rises  sharply  as  the  concentration  decreases.  This 
aberration  of  the  curve  is  unquestionably  due  to  the  presence  of  impurities 
in  the  water  as  a  result  of  which  a  portion  of  the  acid  is  neutralized.  As 
purer  water  is  employed,  this  effect  becomes  less  pronounced  and  ulti- 
mately disappears  entirely.  In  Run  2,  the  point  at  the  lowest  concen- 
tration still  shows  the  effect  to  a  slight  extent.  In  this  run,  resistance 
measurements  were  taken  over  a  period  of  3  hours  after  the  first  addition 
of  acid.  These  data  showed  a  steady  increase  in  resistance.  The  re- 
sistance was  plotted  as  a  function  of  the  time  and  the  curve  was  extra- 
polated to  zero  time.  The  resistance  at  the  end  of  3  hours  was  13668.6, 
while  the  extrapolated  resistance  read  13623.5.  In  order  to  correct  the 
rest  of  the  run  for  this  effect,  the  weight  of  acid  added  was  multiplied  by 


16 


the  ratio  of  these  resistances.  This  run  has  been  corrected  throughout 
in  this  manner,  otherwise  the  second  and  third  points  would  rise  somewhat 
above  the  positions  as  indicated  on  the  plot.  The  fourth  point,  however, 
would  not  be  measurably  affected.  The  remaining  runs  in  quartz  at  low 
concentrations  were  all  carried  out  with  water  of  a  high  degree  of  purity 
and  no  corrections  were  applied. 
26.25 


26.15 


26.05 


2  25.95 
X 

3-25.85 


25  75 


25.65 


J  e=      X 

S          9 


0.0 


60 


1.0       2.0        3.0        4.0        5.0 

Specific  conductance  X  104. 

Fig.  6. — 1/A-CA  plot  for  aqueous  iodic  acid  solu- 
tions at  25°. 

The  points  which  are  most  likely  to  be  in  error  in  this  work  are  those 
at  the  lowest  concentration,  where  the  weight  of  acid  is  small  and  the 
errors  due  to  weighing  have  an  appreciable  influence,  while  the  effect  of 
impurities  in  the  water  is  relatively  great. 

At  high  concentrations,  polarization  effects  may  make  their  appearance. 
With  the  exception  of  the  last  point  in  Run  10,  using  a  glass  cell,  no  resist- 
ance was  measured  below  250  ohms.  The  last  mentioned  effect  may 
therefore  be  considered  negligible.  During  Run  6,  it  was  found  that  a 
slight  crack  had  developed  in  one  of  the  electrode  lead  tubes,  which  per- 
mitted a  trace  of  acid  to  act  on  the  mercury.  This  is  unquestionably  the 
cause  for  the  displacement  of  the  points  obtained  in  this  run.  In  Run  7, 
the  crack  had  been  partially  mended  and  the  results,  as  may  be  observed, 
were  considerably  improved,  but  the  error  is  still  appreciable.  The  cause 
of  the  error  was  so  obvious  and  the  effect  was  so  small  at  the  higher  con- 
centrations that  it  was  thought  unnecessary  to  repeat  the  entire  series 
of  measurements.  With  the  exception  of  the  measurements  just  men- 
tioned, the  conductance  data  in  quartz  cells  are  consistent  within  nearly 
0.01%. 


17 

Relative  Values  of  the  Specific  Conductance  of  0.01  Ar  Potassium  lodate 
and  lodic  Acid  at  18°  and  25°. — In  order  to  derive  the  value  for  the  con- 
ductance of  the  hydrogen  ion  at  25°,  the  limiting  value  of  the  equivalent 
conductance  of  the  iodate  ion  at  that  temperature  must  be  known ;  while, 
to  derive  that  at  18°,  the  limiting  value  of  the  equivalent  conductance  of 
iodic  acid  at  that  temperature  must  be  known.  The  equivalent  conduct- 
ance of  the  iodate  ion  at  18°  has  been  derived  by  Noyes  and  Falk10  from 
the  conductance  measurements  of  Kohlrausch,11  with  potassium  iodate. 
They  assign  to  it  the  value  34.0,  basing  this  on  the  A0  value  98.5  for  po- 
tassium iodate  and  the  value  0.496  for  the  transference  number  of  the 
potassium  ion  in  potassium  chloride  at  18°. 

The  conductance  of  the  iodate  ion  at  25°  has  not  been  determined. 
For  the  purpose  of  determining  the  limiting  value  of  the  conductance 
of  this  ion  at  25  °,  the  specific  conductance  of  potassium  iodate  was  mea- 
sured at  a  concentration  of  0.00149105  N.  This  concentration  is  based 
on  the  density  values  0.99733  and  0.99888,  at  25°  and  18°,  respectively.12 
The  concentrations  and  corresponding  values  of  A  as  measured  are  given 
in  the  following  table. 

TABLE  IX 

CONDUCTANCE  OF  DILUTE  POTASSIUM  IODATE  SOLUTIONS  AT  18°  AND  25° 
Temperature  Concentration  A 

°c. 

25  0.00149105  110.855 

18  0.00149337  95.4318 

Correcting  the  measured  specific  conductance  at  18°  for  the  change  in 
concentration  due  to  temperature  change,  for  the  ratio  of  the  specific 
conductances  at  25°  and  18°  the  value  1.16158  is  obtained. 

In  order  to  check  this  value,  a  smooth  curve  was  drawn  through  Kohl- 
rausch's  values  at  18°  and  a  conductance  value  for  the  concentration 
0.00149105  was  interpolated  and  found  to  be  95.487.  Dividing  this  by 
the  observed  value  of  the  equivalent  conductance  given  in  the  table  above 
for  25  °,  we  have  for  the  ratio  of  the  conductances  at  the  two  temperatures 
the  value  1.16095,  which  agrees  with  the  above-determined  value  within 
0.054%.  Multiplying  98.5,  the  value  of  A0  for  potassium  iodate  at  18°, 
as  determined  by  Noyes  and  Falk,  by  the  ratio  1.16158,  we  obtain  114.42 
for  the  value  of  A0  for  potassium  iodate  at  25  °. 

The  specific  conductance  of  iodic  acid  at  a  concentration  of  0.001  N 
was  determined  at  18°  and  25°  in  Cells  II  and  III.  The  conductance 
at  18°  was  corrected  for  the  concentration  change  due  to  the  temperature 
change  from  25°  to  18°.  The  ratio  of  the  conductances  at  the  same 

10  Noyes  and  Falk,  J.  Am.  Chem.  Soc.,  34,  454  (1912). 

11  Kohlrausch,  Sitzb.  kongl.  preuss.  akad.  Wiss.,  1900,  p.  1002. 

12  Sullivan,  Z.  physik.  Chem.,  28,  525  (1899). 


18 

concentration  in  the  2  cells  was  found  to  be  1.11432  and  1.11410,  or,  in 
the  mean,  1.11421. 

Discussion 

The  Influence  of  Impurities  on  the  Conductance  of  Dilute  Acid  Solu- 
tions.— Reference  has  already  been  made  in  an  earlier  section  of  this  paper 
to  the  influence  of  impurities  on  the  conductance  of  dilute  acid  solution. 
The  influence  of  impurities  is  to  reduce  the  observed  conductance  below 
the  true  value.  As  a  consequence,  the  conductance  values  of  dilute  acid 
solutions  pass  through  a  maximum  value,  instead  of  approaching  a  definite 
limit.  The  weaker  the  acid  and  the  lower  the  concentration,  the  more 
pronounced  is  this  effect. 

The  influence  of  impurities  in  the  water  and  of  the  alkali  due  to  glass 
cells  is  illustrated  in  Fig.  5,  in  which  values  of  the  equivalent  conductance 
A  are  plotted  as  ordinates  and  values  of  the  logarithms  of  the  concentration 
as  abscissas.  Curve  I  represents  the  results  obtained  with  the  quartz  cells 
II  and  IV,  Runs  4,  5,  6  and  7  being  plotted.  Runs  2  and  3  lie  on  the  same 
curve  but  have  been  omitted  for  the  sake  of  clearness  in  the  figure.  Curve 
II  represents  the  results  obtained  with  the  quartz  cell  II,  with  water 
having  a  specific  conductance  of  0.5  X  10  ~6.  Curve  III  represents  the 
results  of  Runs  9  and  10,  with  the  soda-lime  glass  cells  V  and  VI,  the  water 
in  the  case  of  the  dilute  solutions  having  a  specific  conductance  of  0.24 
X  10~6.  Curve  IV  represents  the  results  of  Runs  8  and  10  with  the  same 
soda-lime  glass  cells,  with  water  having  a  specific  conductance  of  0.8 
X  10  ~6.  It  is  seen  that,  with  the  quartz  cells  and  water  having  a  specific 
conductance  of  0.115  X  10  ~6,  or  lower,  a  maximum  does  not  appear  in 
the  conductance  curve  down  to  the  lowest  concentrations  measured. 
Curve  II,  however,  the  measurements  for  which  were  carried  out  with 
water  having  a  specific  conductance  of  0.5  X  10  ~6,  exhibits  a  slight  maxi- 
mum. The  difference  between  Curves  I  and  II  is,  therefore,  due  to  the 
influence  of  impurities  present  in  the  water.  The  pronounced  maximum 
in  Curve  III  is  largely  due  to  the  influence  of  the  alkali  in  the  glass,  since 
in  this  case  the  water  had  a  specific  conductance  of  0.24  X  10 ~6.  This 
curve  lies  much  below  Curve  II,  in  which  case  the  water  had  a  specific 
conductance  of  0.5  X  10  ~6.  The  difference  between  Curve  III  and  Curve 
IV  is  due  to  the  water,  since  the  measurements  were  carried  out  under 
otherwise  comparable  conditions.  It  will  be  observed  that  the  shift  in 
the  curves  due  to  a  change  in  the  specific  conductance  of  the  water  from 
0.24  to  0.8  is  markedly  lower  than  that  due  to  substitution  of  glass  for 
quartz,  which  is  represented  approximately  by  Curves  I  and  III.  With 
the  water  used  in  these  measurements,  the  influence  of  the  alkali  derived 
from  the  cell  is,  therefore,  markedly  greater  than  that  of  the  impurities 
in  water  having  a  specific  conductance  of  0.8  X  10 ~8. 

Water  having  a  specific  conductance  of  0.8  X  10 ~6  is  what  would  be 
ordinarily  termed  "conductivity  water."  It  is  evident  that,  with  water 


19 

of  this  degree  of  purity  in  glass  cells,  it  is  not  possible  to  obtain  reliable 
conductance  measurements  at  concentrations  below  3.X  10  ~3  N.  With 
water  of  a  high  degree  of  purity  in  glass  cells,  the  conductance  measure- 
ments might  be  extended  to  10~3  N.  With  water  of  a  specific  conductance 
0.5  X  10 ~6  in  quartz  cells,  the  influence  of  the  impurities  in  the  water 
becomes  felt  at  a  concentration  of  about  5  X  10 ~4  N.  The  influence  of 
impurities  on  the  conductance  of  dilute  solutions  of  acids  is  brought  out 
more  distinctly  in  plotting  the  reciprocal  of  the  equivalent  conductance 
against  the  specific  conductance,  as  in  Fig.  6.  An  inspection  of  this  figure 
shows  that  the  points  of  Run  1  lie  on  a  curve  which  exhibits  a  pronounced 
minimum  in  the  neighborhood  of  1.3  X  10  ~4  N.  But  even  the  results  of 
Run  2,  which  do  not  exhibit  a  maximum,  diverge  markedly  at  the  lowest 
concentrations.  It  is  evident  that,  if  the  measurements  had  been  carried 
to  lower  concentrations  in  this  case,  a  minimum  would  have  been  found. 
Some  further  evidence  may  be  given  relative  to  the  influence  of  glass 
cells  on  the  results  obtained  with  dilute  acids.  Even  in  the  case  of  the 
Pyrex  glass  cell  III,  the  conductance  as  measured  in  Run  11,  with  water 
having  a  specific  conductance  of  0.53  X  10  ~6,  is  markedly  lower  than  in 
Run  1  with  the  quartz  cell,  with  water  having  approximately  the  same 
specific  conductance.  The  results  with  the  Pyrex  glass  cell  are  indicated 
in  Fig.  5  by  combined  cross  and  circle.  In  order  to  test  out  this  effect 
further,  the  cell  constant  of  this  cell  was  checked  out  against  the  large 
quartz  cell  II  immediately  after  Run  11.  First,  a  solution  of  potassium 
chloride  was  made  up  in  the  large  quartz  cell  and  its  resistance  measured, 
after  which  the  solution  was  transferred  to  the  Pyrex  cell  and  its  resistance 
again  measured.  The  value  obtained  for  the  constant  of  the  Pyrex  cell 
in  this  case  was  8.26792,  which  checks  closely  with  the  value  8.26717 
as  determined  directly  by  means  of  0.1  N  standard  potassium  chloride 
solution.  Then  an  acid  solution,  having  a  concentration  of  approximately 
0.001  A7,  whose  resistance  was  approximately  the  same  as  that  of  the 
potassium  chloride  solution,  was  also  made  up  in  the  quartz  cell,  and  its 
resistance  in  this  cell  compared  with  that  in  the  Pyrex  cell.  The  cell 
constant  obtained  with  the  acid  was  found  to  be  8.28604,  a  difference  of 
0.2%,  by  direct  check  at  the  same  resistance.  This  effect  is  evidently 
due  to  the  action  of  the  alkali  of  the  glass  cell.  The  relative  magnitude 
of  this  effect  will  naturally  depend  upon  the  conditions  under  which  the 
experiments  are  carried  out,  as  well  as  upon  the  dimensions  of  the  cell,  etc. 
The  effect  of  the  glass  is  most  clearly  shown  by  introducing  an  acid  solu- 
tion of  a  concentration  in  the  neighborhood  of  2  X  10~3  N  into  a  glass  cell 
and  thereafter  measuring  the  resistance  as  a  function  of  the  time.  When 
left  in  Cell  V  for  a  period  of  30  minutes,  it  was  found  that  the  resistance 
increased  by  approximately  0.04%.  On  shaking  the  solution,  however, 
a  further  increase  in  resistance  took  place  of  the  same  order  of  magnitude. 


20 

This  effect  is  unquestionably  due  to  the  fact  that  the  concentration  of  the 
impurities  arising  from  the  glass  is  greater  in  the  immediate  neighborhood 
of  the  cell  walls  than  in  the  vicinity  of  the  electrodes.  On  mixing  the 
contents  of  the  cell,  these  impurities  are  distributed  throughout  the  volume 
of  the  solution  with  a  consequent  decrease  of  the  observed  resistance.  It 
may  be  noted  that,  when  the  original  nieasurements  were  carried  out  in 
the  glass  cell,  this  effect  was  not  observed  and  did  not  lead  to  irregularities 
in  the  measured  values,  for  the  reason  that  the  conductance  of  the  solu- 
tions was  measured  at  each  concentration  after  the  same  interval  of  time 
(45  minutes). 

In  the  case  of  solutions  of  weaker  acids,  a  maximum  in  the  conductance 
values  at  concentrations  down  to  2  X  10 "4  N  has  not  been  observed. 
This  is  due  to  the  fact  that  the  maximum  cannot  appear  until  the  decrease 
of  conductance,  due  to  the  influence  of  impurities,  overbalances  the  increase 
due  to  the  increased  ionization  of  the  acids  as  the  concentration  is  de- 
creased. The  weaker  the  acid,  the  greater  the  relative  conductance 
change  for  a  given  concentration  change.  Consequently,  it  is  only  in 
the  case  of  the  strong  acids  that  this  effect  is  observed  at  higher  concen- 
trations. This  doubtless  accounts  for  the  absence  of  such  an  effect  in  the 
measurements  of  Kendall.13  Nevertheless,  in  the  case  of  the  weak  acids, 
the  effects  due  to  impurities  are  present  and  influence  the  observed  con- 
ductance values,  and  consequently  also  the  values  of  A0  obtained  from  such 
nieasurements  by  extrapolation. 

Many  instances  of  this  effect  may  be  found  in  the  literature,  although, 
apparently,  in  few  cases  has  this  effect  been  studied  closely,  while  in  other 
cases  it  has  not  been  correctly  interpreted.  Noyes  and  Kato,14  in  their 
transference  experiments  with  nitric  and  hydrochloric  acids,  found  a 
decrease  of  the  specific  conductance  of  their  0.002  N  acid  solutions  amount- 
ing to  0.1%  on  standing  for  3  hours.  This  is  one  of  the  few  cases  in  which 
this  effect  has  been  measured. 

It  is,  of  course,  obvious  that  this  effect  will  make  itself  felt,  not  only  in 
the  case  of  conductance  measurements,  but  likewise  in  that  of  all  other 
measurements  which  depend  primarily  upon  the  concentration  of  hydrogen 
ions.  So,  for  example,  Beans  and  Oakes15  have  measured  the  concentra- 
tion of  the  hydrogen  ion  in  so-called  pure  water  by  means  of  concentration 
cells.  Their  water  had  a  specific  conductance  of  the  order  0.9  X  10  ~6 
and  the  value  which  they  obtained  for  the  concentration  of  the  hydrogen 
ions  was  about  Vio  that  determined  by  other  methods.  Since  these  mea- 
surements were  carried  out  in  glass  cells,  it  is  clear  that  the  impurities  in 
the  water,  as  well  as  the  alkali  due  to  the  glass  cells,  must  have  had  a  con- 

18  Kendall,  /.  Chem.  Soc.,  101,  1375  (1912). 

14  Noyes  and  Kato,  J.  Am.  Chem.  Soc.,  30,  323  (1908). 

16  Beans  and  Oakes,  ibid.,  42,  2128  (1920). 


21 

siderable  influence  on  the  resulting  value  of  the  hydrogen-ion  concentra- 
tion as  measured. 

Nearly  all  determinations  of  hydrogen-ion  concentrations  have  been 
carried  out  in  glass  cells  with  water  which,  at  best,  has  a  purity  corre- 
sponding with  that  of  what  is  commonly  known  as  "conductivity  water/' 
which  has  a  specific  conductance  of  1  X  10  ~6.  It  is  clear  that  all  such 
measurements  in  which  the  acid  concentration  approaches  10~3  N  must 
be  measurably  in  error,  while  at  concentrations  below  this  value  the  errors 
reach  correspondingly  greater  values.  Without  doubt,  these  effects  had 
an  influence  on  the  electromotive-force  determinations  of  Noyes  and 
Ellis,16  which  led  them  to  state  that  "the  plot  indicates  that  the  value  of 
the  electromotive  force  at  0.000999  N  is  affected  by  a  considerable  error." 
So,  also,  this  effect  undoubtedly  underlies  the  observations  of  Linhart17 
with  the  hydrogen  electrode  at  0.000136  N,  which  led  to  a  decrease  of 
1.5%  in  the  electromotive  force  on  standing  over  a  period  of  126  hours, 
while  in  the  case  of  the  more  concentrated  solutions  an  increase  of  the 
electromotive  force  occurred  under  the  same  conditions. 

It  may  be  safely  assumed  that  all  electromotive-force  measurements 
with  the  acids,  carried  out  at  concentrations  approaching  10~3  N,  are  in 
error  by  varying  amounts  depending  upon  the  length  of  time  during  which 
the  solution  was  left  in  contact  with  the  glass  surfaces,  as  well  as  upon 
other  factors. 

Form  of  the  Conductance  Curve  at  Low  Concentrations. — If  the  law 
of  mass  action  is  approached  as  a  limiting  form,  then,  as  Kraus  and  Bray 
have  pointed  out,18  the  curve  in  which  values  of  I/A  are  plotted  against 
values  of  the  specific  conductance,  becomes  a  straight  line  at  low  con- 
centrations. The  curve  in  Fig.  6  has  been  drawn  to  pass  through  the 
points  within  the  limits  of  experimental  error,  and,  where  the  deviation 
from  a  linear  relation  lies  within  the  limit  of  the  experimental  error,  the 
curve  has  been  continued  as  a  straight  line.  Whether  or  not  the  mass- 
action  law  is  approached  as  a  limiting  case  in  aqueous  solutions  of  strong 
electrolytes,  it  will  be  admitted  that  the  value  of  A0  obtained  by  this 
linear  extrapolation  represents  a  minimum  value.  As  may  be  seen  from 
Fig.  6,  at  a  concentration  below  2  X  10  ~4  A/",  the  points  lie  upon  a  straight 
line  within  the  limits  of  the  experimental  error.  In  carrying  out  the 
extrapolation,  somewhat  greater  weight  has  been  given  to  the  results 
of  Runs  4  and  5  than  on  preceding  runs,  since  in  carrying  out  the  later 
runs  the  details  of  manipulation  had  been  worked  out  more  minutely  and 
the  results  themselves  are  mutually  more  consistent.  The  value  of  A0, 
as  thus  determined  by  extrapolation,  was  found  to  be  389.55.  The  true 

16  Noyes  and  Ellis,  J.  Am.  Chem.  Soc.,  39,  2542  (1917). 

17  Linhart,  ibid.,  41,  1178  (1919). 

18  Kraus  and  Bray,  ibid..  35,  1337  (1913). 


22 

value  of  Ao  cannot  differ  greatly  from  this  value,  for,  at  higher  concen- 
trations, the  curvature  of  the  above  curve  diminishes  as  the  concentration 
decreases.  If,  beginning  at  a  concentration  of  2  X  10  ~4  N,  the  curves 
were  continued  with  the  same  .curvature,  the  resulting  value  of  A0  ob- 
tained would  differ  from  that  found  above  by  only  a  few  tenths  of  a  unit. 

The  value  of  the  mass-action  constant,  corresponding  to  the  straight 
line  drawn  in  Fig.  6,  is  0.0717.  It  is  clear  that  the  value  of  the  ionization 
function  for  iodic  acid  at  a  given  concentration  is  much  greater  than  that 
of  ordinary  salts,  such  as  potassium  chloride,  and  that  the  limiting  value 
approached,  if  such  a  limit  exists,  is  much  greater  than  that  of  salts.19 
Comparing  the  ionization  of  iodic  acid  with  that  of  the  salts,  this  acid  is 
to  be  classed  as  a  strong  electrolyte.  At  low  concentrations,  certainly, 
it  does  not  differ  greatly  in  strength  from  such  strong  acids  as  hydrochloric 
and  nitric  acids. 

The  Conductance  of  the  Hydrogen  Ion  at  25°  and  18°.— From  the 
data  given  in  the  preceding  section,  the  conductance  of  the  hydrogen 
ion  may  be  closely  approximated  at  25°  and  18°.  From  the  ratio  of  the 
conductance  of  10~3  N  potassium  iodate  solutions  at  25°  and  18°  the 
value  114.42  was  obtained  for  the  limiting  value  of  the  equivalent  conduct- 
ance of  potassium  iodate.  Subtracting  the  value  of  74.8  for  the  conduct- 
ance of  the  potassium  ion  at  25  °,20  we  obtain  the  value  39.62  for  the  conduct- 
ance of  the  iodate  ion  at  this  temperature.  This  yields  for  the  conduct- 
ance of  the  hydrogen  ion  at  25°  the  value  349.93.  One  of  the  uncertainties 
underlying  the  value  of  the  conductance  of  the  iodate  ion  given  above  is 
due  to  the  assumption  that  the  ratio  of  the  A0  values  for  potassium  iodate 
at  18°  and  25°  is  equal  to  that  of  the  conductance  values  of  the  same  salt 
at  a  concentration  in  the  neighborhood  of  10  ~3  Ar.  It  is  believed,  however, 
that  the  error  introduced  in  this  way  is  small. 

The  value  thus  obtained  for  the  equivalent  conductance  of  the  hydrogen 
ion  is  nearly  1%  greater  than  that  obtained  by  Kendall13  from  the  con- 
ductance of  solutions  of  weak  acids.  Whatever  other  errors  may  have 
affected  Kendall's  results,  it  is  certainly  true  that  the  influence  of  im- 
purities in  the  water,  as  well  as  the  influence  of  alkali  from  the  cell  walls, 
might  be  expected  to  lead  to  low  values  of  the  above  order  of  magnitude. 

In  Section  V,  the  ratio  for  the  equivalent  conductance  of  iodic  acid  at 
a  concentration  of  10~3  N  at  25°  and  18°  was  found  to  be  1.11421.  Di- 
viding 389.55,  the  value  of  A0  for  iodic  acid  at  25°,  by  this  ratio,  we  obtain 
the  value  349.62  for  A0  of  this  acid  at  18°.  Subtracting  34.0,  the  con- 
ductance of  the  iodate  ion,  as  derived  by  Noyes  and  Falk,  we  obtain  for 

19  The  maximum  value  which  might  be  assigned  to  the  limit  approached  by  the  mass 
action  function  in  the  case  of  KC1  is  0.02  according  to  Washburn  and  Weiland  [/.  Am. 
Chem.  Soc.,  40,  146  (1918)]. 

20  This  value  is  that  derived  by  Noyes  and  Falk  (Ref.  10),  assuming  the  value  0.497 
for  the  transference  number  of  the  potassium  ion  in  potassium  chloride  at  25°. 


23 

the  conductance  of  the  hydrogen  ion  at  18°  the  value  315.62.  This  value 
is  slightly  higher  than  that  derived  by  Noyes  and  Falk.  It  would  appear, 
however,  that  this  value  would  necessarily  represent  a  lower  limit  in  view 
of  the  method  of  extrapolation  employed  in  determining  the  value  of  A0. 

Summary 

1.  The  apparatus  employed  and  the  precautions  observed  in  carrying 
out  conductance  measurements  with  iodic  acid  at  concentrations  down 
to  5  X  10  ~5  N  are  described.     Measurements  were  carried  out  in  glass 
and  in  quartz  cells  and  with  water  of  various  degrees  of  purity  in  order  to 
determine  the  influence  of  impurities  on  the  conductance  of  acid  solutions 
at  low  concentration. 

2.  It  was  found  that  the  conductance  curve  exhibits  a  maximum  due 
to  impurities  with  water  having  a  specific  conductance  above  0.1  to  0.2 
X  10  ~6.     The  influence  of  alkali  derived  from  glass  cells  is,  if  anything, 
greater  than  that  of  the  impurities  present  in  the  water  having  a  specific 
conductance  of  0.8  X  10 ~6.     Conductance  measurements  with  iodic  acid 
in  quartz  cells  with  water  having  a  specific  conductance  of  0.1  X  10  ~6 
were  carried  to  concentrations  as  low  as5X10~5Af  with  a  relative  pre- 
cision of  a  few  hundred ths  of  1%.     Extrapolating  on  the  assumption  that 
the  mass-action  law  is  approached  as  a  limiting  form  at  low  concentrations, 
389.55  is  found  for  the  value  of  Ao  of  iodic  acid  at  25°,  which  may  be  ac- 
cepted as  a  lower  limit  to  the  possible  value  of  this  constant.     The  mass- 
action  constant  corresponding  to  the  extrapolation  has  a  value  of  0.0717. 
Iodic  acid  is  thus  a  much  stronger  electrolyte  than  potassium  chloride. 

3.  The  conductance  of  the  iodate  ion  was  evaluated  at  25°  on  the 
assumption  that  the  A0  value  between  18°  and  25°  changes  in  the  same 
proportion  as  the  conductance  of  a  0.0015  N  solution  of  the  acid.     The 
value  39.62  was  thus  found  for  the  conductance  of  the  iodate  ion.     For 
the  conductance  of  the    hydrogen  ion  at  25°  the  value  349.93  results. 
Assuming  that  the  limiting  value  of  the  equivalent  conductance  of  iodic 
acid  between  18°  and  25°  varies  in  the  same  ratio  as  that  of  a  0.001  N 
solution  of  the  same  acid,  349.62  is  found  for  the  value  cf  A0  of  iodic  acid 
at  18°.     Assuming  the  value  34.0  for  the  equivalent  conductance  of  the 
iodate  ion  at  18°,  there  is  obtained  the  value  315.62  for  the  conductance 
of  the  hydrogen  ion  at  18°. 


507f 


P3 


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